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Dr. Vatsala Soni PGGC Sector -11 Chandigarh
The s-Block Elements Dr. Vatsala Soni PGGC Sector Chandigarh
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Members of the s-Block Elements
Li Be Na K Rb Cs Fr Mg Ca Sr Ra Ba IA IIA IA Alkali metals IIA Alkaline Earth metals
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Chapter summary Characteristic properties of the s-block elements
Variation in properties of the s-block elements Variation in properties of the s-block compounds Uses of compounds of the s-block elements
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Characteristic properties of s-block elements
Metallic character Low electronegativity Basic oxides, hydroxides Ionic bond with fixed oxidation states Characteristic flame colours Weak tendency to from complex
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Metallic character High tendency to lose e- to form positive ions
Metallic character increases down both groups
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Electronegativity Low nuclear attraction for outer electrons
Highly electropositive Small electronegativity Group I Group II Li Be Na Mg 1.2 K Ca Rb Sr Cs Ba Fr Ra
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Basic oxides, hydroxides
Li2O LiOH Na2O, Na2O2 NaOH K2O2, KO2 KOH Rb2O2, RbO2 RbOH Cs2O2, CsO2 CsOH Oxide Hydroxides BeO Be(OH)2 MgO Mg(OH)2 CaO Ca(OH)2 SrO Sr(OH)2 BaO, Ba2O2 Ba(OH)2
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Oxides, Peroxide, Superoxide
Reaction with water: Oxide: O2- + H2O 2OH- Peroxide: O H2O H2O2 + 2OH- Superoxide: 2O2- + 2H2O 2OH- + H2O2 + O2 :O:.O: .. .. Super oxide :O:O: .. .. Peroxide ion Li does not form peroxide or super oxide Li2O2 Li2O + ½ O2
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Hydroxides Group I hydroxides Li Na K Rb Cs
All are soluble, base strength increase. Group II hydroxide Be Mg Ca Sr Ba Solubility increase, from Amphoteric to basic, base strength increase
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Predominantly ionic with fixed oxidation state
Group I: Most electropositive metals. Low first I.E. and extremely high second I.E. Form predominantly ionic compounds with non-metals by losing one electron. Fixed oxidation state of +1. Group II: Electropositive metals. Low first and second I.E. but very high third I.E Have a fixed oxidation state of +2. Be and Mg compounds possess some degree of covalent character.
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Characteristic flame colours
Na+ Cl- (g) Na (g) + Cl (g) Na(g) Na* (g) [Ne]3s [Ne]3p1 Na*(g) Na(g) + h (589nm, yellow)
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Flame test Li deep red Ca brick red Na yellow Sr blood red K lilac
Rb bluish red Cs blue Ca brick red Sr blood red Ba apple green HCl(aq) sample
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Weak tendency to form complex
Complex formation is a common feature of d-block element. e.g. Co(NH3)63+ Co :NH3 H3N: s-block metal ions have no low energy vacant orbital available for bonding with lone pairs of surrounding ligands, they rarely form complexes.
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Variation in properties of elements
Atomic radii Ionization enthalpies Hydration enthalpies Melting points Reactions with oxygen, water, hydrogen and chlorine
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Atomic radii (nm) Li 0.152 Be 0.112 Na 0.186 Mg 0.160 K 0.231 Ca 0.197
Rb 0.244 Sr 0.215 Cs 0.262 Ba 0.217 Fr 0.270 Ra 0.220 Fr Ra Li Be
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Ionization Enthapy Group I 1st I.E. 2nd I.E. Li 519 7300 Na 494 4560 K
418 3070 Rb 402 2370 Cs 376 2420 Group I 1st I.E. 2nd I.E. 3rd I.E. Be 900 1760 14800 Mg 736 1450 7740 Ca 590 1150 4940 Sr 548 1060 4120 Ba 502 966 3390
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Ionization Enthalpy 1st I.E. Be+ Li Na 2nd IE K Rb Ca+ Cs Ba+ Be Ca Ba
300 400 500 600 500 1000 1500 2000 Be Ca Ba Be+ Ca+ Ba+ 1st IE 2nd IE
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Ionization Enthalpy Group I
Have generally low 1st I.E. as it is well shielded from the nucleus by inner shells. 2. Removal of a 2nd electron is much more difficult because it involves the removal of inner shell electron. 3. I.E. decreases as the group is descended. As atomic radius increases, the outer e is further away from the well-shielded nucleus.
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Ionization Enthalpy Group II Have low 1st and 2nd IE.
Removal of the 3rd electron is much more difficult as it involves the loss of an inner shell electron. IE decrease as the group is descended. IE of the group II is generally higher than group I.
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Hydration Enthalpy M+(g) + aqueous M+(aq) + heat M+ -600 -300
Li+ Na+ K+ Rb+ Cs+ M+
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Hydration Enthalpy Be2+ Mg2+ Ca2+ Sr2+ Ba2+ -2250 -2000 -1750 -1500
-600 -300 Li+ Na+ K+ Rb+ Cs+
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Hydration Enthalpy General trends:
On going down both groups, hydration enthalpy decreases. (As the ions get larger, the charge density of the ions decreases, the electrostatic attraction between ions and water molecules gets smaller.) Group 2 ions have hydration enthalpies higher than group 1. ( Group 2 cations are doubly charged and have smaller sizes)
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Variation in Melting Points
250 500 750 1000 1250 Be Mg Ca Sr Ba Li Na K Rb Cs
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Variation in Melting Points
Strength of metallic bond depends on: Ionic radius Number of e- contributed to the electron sea per atom Crystal lattice structure Note: The exceptionally high m.p. of calcium is due to contribution of d-orbital participation of metallic bonding.
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Variation in Melting Points
Group I Structure Group II Li B.C.C. Be H.C.P. Na Mg K Ca C.C.P. Rb Sr Cs Ba
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Reactions with oxygen S-block elements are strong reducing agents.
Their reducing power increases down both groups. (As the atomic size increases, it becomes easier to remove the outermost electron) S-block elements reacts readily with oxygen. Except Be and Mg, they have to be stored under liquid paraffin to prevent contact with the atmosphere.
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Reactions with oxygen Normal Oxide Peroxide Superoxide Structure
Formed by Li and Group II Na and Ba K, Rb, Cs .. 2- :O: .. :O-O: .. .. :O:.O: .. ..
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Reaction with water M(s) M+(aq) + e- H2O(l) + e- OH-(aq) + ½ H2(g)
Li volt Na K Rb Cs Be volt Mg -2.38 Ca Sr Ba Energetic vs. Kinetic Factor
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Reaction with hydrogen
All the s-block elements except Be react directly with hydrogen. 2Na(s) + H2(g) 2NaH(s) Ca(s) + H2(g) CaH2(s) The reactivity increases down the group. Only BeH2 and MgH2 are covalent, others are ionic.
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Reaction with chlorine
All the s-block metals react directly with chlorine to produce chloride. All group I chlorides are ionic. BeCl2 is essentially covalent, with comparatively low m.p. The lower members in group II form essentially ionic chlorides, with Mg having intermediate properties.
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Although lithium has highly negative Eo, it only
reacts slowly with water. This illustrates the importance of the role of kinetic factors in determining the rate of a chemical reaction. Lithium has a higher m.p., this increases the activation energy required for dissolution in aqueous solution. It does not melt during the reaction as Na and K do, and thus it has a smaller area of contact with water.
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Variation in properties of the compounds
Reactions of oxides and hydroxides Reactions of chlorides Reactions of hydrides Relative thermal stability of carbonates and hydroxides Relative solubility of sulphate(VI) and hydroxde
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Reactions of oxides and hydroxides
All group I oxides reacts with water to form hydroxides Oxide: O2- + H2O 2OH- Peroxide: O H2O H2O2 + 2OH- Superoxide: 2O2- + 2H2O 2OH- + H2O2 + O2 All group I oxides/hydroxides are basic and the basicity increases down the group.
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Reactions of oxides and hydroxides
Group II oxides/hydroxides are generally less basic than Group I. Beryllium oxide/hydroxide are amphoteric.
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Reactions of chlorides
All group I chlorides are ionic and readily soluble in water. No hydrolysis occurs. Group II chlorides show some degree of covalent character. Beryllium chloride is covalent and hydrolysis to form Be(OH)2(s) and HCl(aq). Magnesium chloride is intermediate, it dissolves and hydrolysis slightly. Other group II chlorides just dissolve without hydrolysis.
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Reactions of hydrides They all react readily with water to give the
metal hydroxide and hydrogen due to the strong basic property of the hydride ion, H:- H:-(s)+ H2O(l) H2(g)+ OH-(aq) Hydride ions are also good reducing agent. They can be used to prepare complex hydrides such as LiAlH4 and NaBH4 which are used to reduce C=O in organic chemistry.
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Thermal Stability Thermal stability refers to decomposition of the
compound on heating. Increased thermal stability means a higher temperature is needed to decompose the compound.
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Thermal Stability of carbonates
Li2CO3 Li2O + CO2 ( at 700oC) All other group I carbonates are stable at ~800oC BeCO3 BeO + CO2 ( at 100oC) MgCO3 MgO + CO2 ( at 540oC) CaCO3 CaO + CO2 ( at 900oC) SrCO3 SrO + CO ( at 1290oC) BaCO3 BaO + CO2 ( at 1360oC)
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Thermal Stability of hydroxides
All group I hydroxides are stable except LiOH at Bunsen temperature. Be(OH)2(s) BeO(s) + H2O(g) H = +54 kJ/mol Mg(OH)2(s) MgO(s) + H2O(g) H = +81 kJ/mol Ca(OH)2(s) CaO(s) + H2O(g) H = +109 kJ/mol Sr(OH)2(s) SrO(s) + H2O(g) H = +127 kJ/mol Ba(OH)2(s) BaO(s) + H2O(g) H = +146 kJ/mol
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Thermal stability Carbonates and hydroxides of Group I metals
are as a whole more stable than those of Group II. Thermal stability increases on descending the group. Lithium often follow the pattern of Group II rather than Group I. This is an example of the diagonal relationship.
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Explanation of Thermal Stability
Charge of the ions Size of the ions Compounds are more stable if the charge increases and size decreases. For compounds with large polarizable anions, thermal stability is affected by the polarizing power of the cations.
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Explanation of Thermal Stability
- + Decreasing polarizing power + Increasing stability - + -
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Explanation of Thermal Stability
Mg2+ C O O:- - :O Mg2+ O2- + CO2 Mg2+ Mg2+ O2- + H2O -:O H
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Explanation of Thermal Stability
MgCO3 MgO MgO BaO BaCO3 BaO
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Relative solubility of Group II hydroxides
Compound Solubility / mol per 100g water Mg(OH)2 0.020 x 10-3 Ca(OH)2 1.5 x 10-3 Sr(OH)2 3.4 x 10-3 Ba(OH)2 15 x 10-3 Solubility of hydroxides increases down the group.
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Solubility of Group II sulphates
Compound Solubility / mol per 100g water MgSO4 3600 x 10-4 CaSO4 11 x 10-4 SrSO4 0.62 x 10-4 BaSO4 0.009 x 10-4 Solubility of sulphates increases up the group.
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Explanation of solubility
MX(s) aqueous H solution M+(aq) + X-(aq) M+(g) + X-(g) H hydration -H lattice H solution -H lattice H hydration = +
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Explanation of solubility
Group I compounds are more soluble than Group II because the metal ions have smaller charges and larger sizes. H lattice is smaller, and H solution is more exothermic. H solution -H lattice H hydration = +
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Explanation of solubility
H solution -H lattice H hydration = + For Group II sulphates, the cations are much smaller than the anions. The changing in size of cations does not cause a significant change in H lattice (proportional to 1/(r+ + r-). However, the changing in size of cations does cause H hydration (proportional to 1/r+ and 1/r-) to become less exothermic, and the solubility decreases when descending the Group.
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SO42- MgSO4 SrSO4
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Explanation of solubility
H solution -H lattice H hydration = + For the smaller size anions, OH-. Down the Group, less enthalpy is required to break the lattice as the size of cation increases. However the change in H solution is comparatively smaller due to the large value of 1/r- . As a result, H solution becomes more exothermic and the solubility increases down the Group.
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Mg(OH)2 Sr(OH)2
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Uses of s-block compounds
Sodium carbonate Manufacture of glass Water softening Paper industry Sodium hydrocarbonate Baking powder Soft drink
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Uses of s-block compounds
Sodium hydroxide Manufacture of soaps, dyes, paper and drugs To make rayon and important chemicals Magnesium hydroxide Milk of magnesia, an antacid Calcium hydroxide To neutralize acids in waste water treatment Strontium compound Fireworks, persistent intense red flame
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Thank You
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